Purpose: The purpose of this lab is to determine the molar mass of solid

tartaric acid. To do this, a solution of sodium hydroxide will first be standardized against a weighed sample of potassium hydrogen phthalate (KHP), an acid whose mass and composition is consistent over time. The sodium hydroxide will then be used to titrate a solution of tartaric acid so that the molar mass of tartaric acid can be accurately calculated.

Background: An acid/base titration is performed by carefully adding one solution

from a buret, the titrant, to another substance in a flask until all of the substance in the flask has reacted. To show when the titration is complete, a dye called an indicator is added to the solution to show when the end point, or the color change indicated when chemical equilibrium between moles of the acid and moles of the base is reached. In the titration performed in this experiment, the reactant NaOH must be standardized against a primary standard such as KHP because it is a compound that changes chemical composition in storage. Solid NaOH absorbs water vapor so that the mass of pellets of NaOH is partly NaOH and partly water. Thus, it is impossible to know the exact molarity of any solution of NaOH without standardizing it by titration because of the variable water content that will distort and invalidate a mass measurement. A heat-dried sample of KHP can be relied on to be 100% KHP; thus, an accurately determined mass of KHP will reveal an accurate number of moles of sodium hydroxide.

Materials Potassium hydrogen

phthalate (KHP), 204. g/mol

NaOH solution (approx. 0.100M)

Tartaric acid Phenolphthalein (Phth)

indicator

Buret and clamp 250ml Erlenmeyer flask Graduated cylinder Distilled water bottle Analytical balance Magnetic stirrers Bottle container

Procedure:Part 1: NaOH and KHP

1) Obtain a clean buret washed with soap and rinsed with distilled water.

2) Calculate mass of NaOH in a 0.1M solution of NaOH, and obtain the calculated mass using weigh boats. Be sure to “tare” the balance with only the weigh boat on it before adding the NaOH.

3) Put massed NaOH in a Florence flask, and add water to the 1L mark.4) Place a magnet into the flask, and the flask on the stirrer. Turn on

the stirrer and dissolve the NaOH.5) Pour the solution into a bottle container.

6) Measure out 0.5g-1.0g of KHP, and record the exact mass. Put the KHP in an Erlenmeyer flask, add about 100mL of distilled water, and dissolve it using the magnet stirrers.

7) Add 2-3 drops of Phenolphthalein indicator into the Erlenmeyer flask.

8) Rinse the buret with a small portion of the NaOH solution. Clamp it so that the buret will drip into the Erlenmeyer flask. Then, using a funnel, fill the buret with the solution to the “0” mark, making sure not to let solution spill over the top.

9) Calculate the amount of NaOH theoretically needed to reach the equivalence point, and add 10mL less than that from the buret to the flask.

10)Open the buret slightly so that NaOH will drip slowly.11)When the color of the solution in the flask turns slightly pink, tighten

the buret, calculate the amount of NaOH used, and record the data.12)Calculate the molarity and label bottle container containing NaOH

solution.13)Repeat titration twice.

Part II: NaOH and Tartaric Acid14) Measure 0.1g-0.2g of tartaric acid, and record exact mass.15) Add about 100mL of water, the measured tartaric acid, 2-3

drops of Phth indicator, and a magnet to a clean Erlenmeyer flask.16) Dissolve the acid. 17) Record the initial volume of NaOH in your buret. Titrate the tartaric

acid solution with the standardized NaOH. Record the final volume of NaOH. Calculate and total volume of NaOH used in the titration.

18) Calculate moles of tartaric acid in your sample as well as the molar mass of tartaric acid in g/mol.

19) Repeat titration until results are within 10% of each other.

Conclusion:The molar mass of solid tartaric acid is 150.0g/mol. This mass

was found by first multiplying the volume of NaOH required to reach the end point (avg. of trials was 21.34mL) and the molarity of the NaOH solution (calculated by dividing moles of KHP by volume of NaOH needed to complete the titration) to calculate moles of NaOH (0.0974M). Since the molar ratio of tartaric acid to NaOH is 1:2, dividing the calculated moles of NaOH by 2 yielded the moles of tartaric acid in the sample (avg. of trials was 0.00104). Lastly, dividing the measured mass of tartaric acid (0.155g) by the calculated moles of acid yielded the molar mass of tartaric acid. The average of the three trials showed the molar mass of tartaric acid to be 150.0g/mol. The calculations used in this experiment largely relied on stoichiometry and the definition of molarity (moles divided by liters) to allow the manipulation of data for the desired information – molar mass of tartaric acid. The chemical process that occurred during the titration of NaOH and tartaric acid involved hydroxide (OH-) ions reacting with the hydrogen (H+) ions floating in the solution of dissolved H2C4H4O6 in the beaker. So, for every mole of acid 2 moles of hydrogen ions reacted with the hydroxide ions dissolved in the solution, as tartaric acid is a diprotic acid. When the equivalence point was reached, signifying chemical equilibrium between moles of H2C4H4O6 acid and moles of the NaOH base, the end point, or the pink color change, indicated that the titration was complete.